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Chapter 10

Chapter 10. Energy. 10.1 The Nature of Energy. Energy: the ability to do work or produce heat Potential energy (store energy): energy due to position or composition Kinetic energy (motion energy): energy due to motion of the object and depends on the mass of the object and its velocity

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Chapter 10

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  1. Chapter 10 Energy

  2. 10.1 The Nature of Energy • Energy: the ability to do work or produce heat • Potential energy (store energy): energy due to position or composition • Kinetic energy (motion energy): energy due to motion of the object and depends on the mass of the object and its velocity • KE = ½ (mv2) • Law of conservation of energy: that energy can be converted from one from to another but can be neither created or destroyed. • Energy of the universe is constant

  3. The nature of energy • Work: force acting over a distance • W = F/d • State function: property of the system that changes independently of its pathway • Which of the following are state functions, and which are not • The temperature of an ice cube • The volume of an aerosol can • The amount of time required for a 10 mi bike ride

  4. 10.2 Temperature and Heat • Temperature: is a measure of the random motions of the components of a substance • E.g H2O molecules move rapidly in hot water than in cold water • Heat: a flow of energy due to a temperature difference • Tfinal = average temp from mixing (hot & cold temp)

  5. 10.3 Exothermic and Endothermic Process • System – everything we focus on in experiment • Surroundings– everything other the system • exothermic (energy flows out of system to surrounding (via heat) • endothermic ( energy flows into system from surrounding (via heat)

  6. Examples • Identify whether these process are exothermic or endothermic • Your hand gets cold when you touch ice • The ice melts when you touch it • Propane is burning in a propane torch • Two chemicals mixing in a beaker give off heat

  7. 10.4 Thermodynamics • Is the study of energy. • First law of thermodynamics: the energy of the universe is constant • Internal energy – energy of the system • ∆E = q x w • ∆ => change in the function • q => represents heat • W => represents work

  8. 10.5 Measuring Energy changes • calorie: the amount of energy (heat) required to raise the temperature of one gram of water by 1oC • 1Calorie = 1000 calories • Joule (J) – SI unit • 1 calories = 4.184 joules

  9. Converting Calories to Joules • Express 60.1 cal of energy in units of joules • How many calories of energy corresponds to 28.4 J?

  10. Specific heat • The amount of energy required to change the temperature of one gram of a substance by 1oC • Denoted as s • Heat required = specific heat x mass x change in temp • q = s m ∆T

  11. Calculating Energy Requirements • Determine the amount of energy (heat) in joules required to raise the temperature of 7.40 g water from 29.0oC to 46.0oC • A 5.63 g sample of solid gold is heated from 21oC to 32oC. How much energy in Joules and calories is required? • A sample of gold requires 3.1 J of energy to change its temperature from 19oC to 27oC. What is the mass of this sample of gold • A 55.0 g aluminum block initially at 27.5oC absorbs 725 J of heat. What is the final temperature?

  12. 10.6 Enthalpy • Heat of reaction (∆H) = the amount of heat that being absorbed or release by a given reaction • ∆Hp = heat

  13. Enthalpy • When 1 mol of methane (CH4) is burned at constant pressure, 890 kJ of energy is released at heat. Calculate ∆H for a process in which a 5.8 g sample of methane is burned at constant pressure • The reaction that occurs in the heat packs used to treat sports injuries is 4Fe(s) + 3 O2(g)  2Fe2O3(s) ∆H= -1652 kJ How much heat is released when 1.00 g of Fe(s) is reacted with excess O2(g)

  14. Entropy • Entropy (S) – measure of disorder or randomness • Randomness increase -> S increases • Sgas> Sion >Smolecule>Ssolid • Ssolid = 0 • E.gH2O(s)  H2O (l) S = + H2O (l)  H2O (g) S = + H2O (g)  H2O (l) S = - • 2nd law of thermodynamic: the entropy of the universe is always increasing

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